Sciencemadness Discussion Board - Why is my dinitrogen tetroxide green?

Why is my dinitrogen tetroxide green?

low.safety.standards - 24-9-2018 at 11:16

Hello people,

I liquefied NO2 with a cold trap and got a deep green color in my N2O4 (instead of deep brown/red). It's the second time I generate it from different sources with different glassware and it turns green. Any ideas?

WGTR - 24-9-2018 at 11:27

Because it has some N₂O₃ dissolved in it. Pure N₂O₃ would be dark blue. How did you generate it?

low.safety.standards - 24-9-2018 at 11:50

Quote: Originally posted by WGTR

Because it has some N₂O₃ dissolved in it. Pure N₂O₃ would be dark blue. How did you generate it?

Ohhh I see... This time it was Copper + RFNA, and Sulfure + RFNA last time.

WGTR - 24-9-2018 at 13:17

In practice, some of the produced NO₂ can react with the water that is formed to produce more nitric and nitrous acids. Invariably some NO is produced, and this combines with NO₂ to form N₂O₃ on condensation. When I react copper with some nitric acid in a loosely stoppered test tube, at first brown gas is produced, then over time the gas color turns clear. Once the stopper is removed briefly, the gas turns brown again as oxygen enters the tube.

If you cool down your condensed product a bit below 0°C, then you can very slowly and carefully let drops of water into the test tube. Don't let them drip into the liquid, but only run down the insides of the test tube, so that no mixing occurs. The water will react with the N₂O₄, creating nitric acid around 68% concentration, with possibly some dissolved N₂O₄. This layer will float on top. On the bottom, intensely blue N₂O₃ will form, like in this post: https://www.sciencemadness.org/whisper/viewthread.php?tid=84...

You could probably take the N₂O₃ and make nitrite from it, but I've never tried it.

low.safety.standards - 24-9-2018 at 22:18

Quote: Originally posted by WGTR

If you cool down your condensed product a bit below 0°C, then you can very slowly and carefully let drops of water into the test tube.

I need it as dry as possible. The purpose is to use as oxidizer for a hypergolic mixture with turpentine. I saw in Purification of Laboratory Chemicals that I can bubble oxygen at 0° to oxidize the N2O3 to N2O4, I assume I have low water content to be getting this color, right?

WGTR - 25-9-2018 at 03:09

You may have a few percent water, but I don’t know how dry you really need it to be. If you had a lot of water, I think you’d see a second layer floating on top. If you want it very dry, you could add water to your fuming acid to dilute it, and then react that with your copper. The primary product would then be nitric oxide, which could be dried by passing it through silica gel. After the drying process, it could then be mixed with oxygen to allow it to oxidize to nitrogen dioxide, and then condensed. With this idea, keep in mind that catalytic trace amounts of water are required for the oxidation to occur, so if the reaction is completely anhydrous you’ll end up with nitric oxide only. In practice, I don’t think you’d end up with conditions that dry to worry about that.

low.safety.standards - 25-9-2018 at 11:21

Quote: Originally posted by WGTR

You may have a few percent water, but I don’t know how dry you really need it to be. If you had a lot of water, I think you’d see a second layer floating on top. If you want it very dry, you could add water to your fuming acid to dilute it, and then react that with your copper.

That's what I did! I tried exactly what you described, the N2O4 was made from 90%+ fuming nitric acid with copper (sulfur previously), the gas was passed through silica cat liter and nothing came out, silica gel absorbs N02 quite efficiently!

fusso - 25-9-2018 at 11:30

Quote: Originally posted by low.safety.standards

Why would you use S to make NO2? S may not be oxidized to H2SO4 completely and some SO2 may form, but I'm not sure about this. Probably decomposing nitrate salts may yield purer NO2/N2O4.

By the way why's this in EM instead of gen chem?

[Edited on 25/09/18 by fusso]

walruslover69 - 25-9-2018 at 11:31

Calcium chloride, would probably make a good desiccant.

woelen - 25-9-2018 at 12:07

No, CaCl2 is a very bad drying agent in this case. I know from experience. It reacts with NO2, the reaction products being mainly ONCl (nitrosyl chloride) and nitrate ion. You also may get some Cl2.

Making pure liquid NO2/N2O4 is very difficult in a home setting. I tried several times but never succeeded. I want a nice sample for display purposes, but I always end up with a dark green/blue liquid. Oxidizing it with O2 might be an option, but with the bubbling O2 you lose most of your volatile liquid as gas. If you have very good cooling, then it may work, but I do not have the equipment for that.

low.safety.standards - 25-9-2018 at 12:44

Quote: Originally posted by fusso

Why would you use S to make NO2? S may not be oxidized to H2SO4 completely and some SO2 may form, but I'm not sure about this. Probably decomposing nitrate salts may yield purer NO2/N2O4.

Fuming HNO3 + S releases lots of NO2! much faster than with Cu, and I wanted to try to convert the H2SO3 to H2SO4 later.

Quote: Originally posted by fusso

By the way why's this in EM instead of gen chem?

I thought people around this section could give me better insights, I'm making "dry as possible" N2O4 for rocket fuel oxidizer.

Quote: Originally posted by woelen

Making pure liquid NO2/N2O4 is very difficult in a home setting. I tried several times but never succeeded. I want a nice sample for display purposes, but I always end up with a dark green/blue liquid. Oxidizing it with O2 might be an option, but with the bubbling O2 you lose most of your volatile liquid as gas. If you have very good cooling, then it may work, but I do not have the equipment for that.

I have a 15 amp peltier I could use to make a cold finger for a methanol bath, I'm going to try and assemble something. I also wanted a pure NO2 ampoule to demonstrate the equilibrium between NO2<->N2O4 at different temperatures. Pic shows that near 0º the red gas tint is greatly reduced, I'll try bubbling in these conditions.

WGTR - 25-9-2018 at 12:59

Quote: Originally posted by low.safety.standards

Ahhh, but you forgot to dilute the acid with water first. With dilute nitric acid you primarily get nitric oxide instead of nitrogen dioxide with concentrated acid. Maybe I can demonstrate it if I can find some time; it shouldn’t take too long.

As an FYI, you can recover the nitrogen dioxide from silica gel by heating it.

low.safety.standards - 25-9-2018 at 13:05

Quote: Originally posted by WGTR

That would explain why I got such a deep green/blue tint, even small droplets are blueish.

Thanks for the tip, I haven't opened the dryer yet to see in what state the silica is in. I know a lot got into it, when I didn't see anything coming through I generated a lot more NO₂, it captured everything.

Quote: Originally posted by fusso

Probably decomposing nitrate salts may yield purer NO2/N2O4.[Edited on 25/09/18 by fusso]

The Handbook of Preparative Inorganic Chemistry says to generate NO₂ from lead nitrate dried in an oven for several days. P₂O₅ is used to dry it further but there are several other steps and a complicated apparatus with cold traps and a furnace. It mentions a path for very pure NO₂ with N₂O₅ but both need phosphorus. Sadly doesn't mention how much you get for each procedure. Possibly generating it from a dried salt alone may be enough for my purpose. I'll try bubbling O₂ and determining the water content of what I have now first.

[Edited on 25-9-2018 by low.safety.standards]

WGTR - 25-9-2018 at 13:17

If you use clear, non-indicating silica gel, it will turn yellowish-brown as it saturates with NO₂. NO by itself will not adsorb onto silica in any appreciable amount. If oxygen is present, then silica will catalyze the oxidation of NO to NO₂, which will then adsorb onto the silica, provided that the silica is dry, and not already saturated with NO₂. NO₂-saturated silica should also catalyze the oxidation of NO, but this is something that I'd like to verify.

[Edited on 9-25-2018 by WGTR]

WGTR - 25-9-2018 at 17:17

Here, I whipped up a quick experiment and took some pictures.

In the picture of the overall setup, this picture was taken after the experiment was completed. I just put everything back together to show how things were assembled.

The test tube on the left contained ≈50% nitric acid and copper wire. The gasses coming from the test tube were bubbled through water (large 1L bottle, almost completely full of water to minimize head space) to remove any residual NO₂, and convert part of it to NO, and also to remove any acid mist. Gasses leaving this bottle were passed through a short train of indicating silica gel. The other end of this tube was pushed tightly through a sealed Ziploc bag. The Erlenmeyer flask isn't part of the experiment.

Even though I took care to exclude oxygen, the silica gel almost immediately turned dark brown. It's possible that the gel itself contained oxygen. The bag filled with a gas that was only vaguely yellow, but filled just the same. After filling the bag half-full with gas, the silica gel/plastic bag assembly was removed, and then a slight excess of dry oxygen (from an oxygen cylinder used for welding) was passed through the silica gel into the bag. As the oxygen flowed into the bag, the gas immediately turned dark brown. Also, the oxygen served to flush NO₂ from the silica gel, changing its color to a light green color.

In my experience, once NO₂ has been adsorbed onto blue indicating silica gel, it will never be blue again, no matter how much heating is applied. It will always be a light green color, even after the NO₂ has been expelled.

After waiting a few minutes, the bag was placed over an open dewar containing liquid nitrogen. One corner of the bag was close to the boiling nitrogen, but was not submerged into it. It was about 1" away from the surface. A freezing salt mixture would probably work just as well, if the bag was submerged a bit, and more time was allowed for condensation.

After waiting for the N₂O₄ to freeze in the corner of the bag, the bag was removed, and allowed to thaw. A straw-colored liquid collected in the bottom corner (picture is rotated).

At no time did I observe any blue or green coloration, in the solid or liquid phase. The frozen material was white, and the liquid was a straw color as seen above.

After returning to the lab 10 minutes later, all of the liquid had evaporated except for a handful of very, very, tiny droplets. I'm not sure what those are. Perhaps it's from some residual moisture.

fusso - 26-9-2018 at 06:04

Quote: Originally posted by WGTR

In my experience, once NO₂ has been adsorbed onto blue indicating silica gel, it will never be blue again, no matter how much heating is applied. It will always be a light green color, even after the NO₂ has been expelled.

Does NO(2) react with/complex to the metal in H2O indicator?

low.safety.standards - 26-9-2018 at 11:04

I'll follow that as I lost all my N2O3 overnight, it ate through what was supposed to be a Teflon cap on that test tube It was stored. Luckly I store high chemicals underwater @home so the gasses were mostly absorbed.

[Edited on 26-9-2018 by low.safety.standards]

DBX Labs - 18-4-2021 at 13:04

A few days back I fed some NO2/NO from a nitrite / H2SO4 decomposition into ampoules cooled down to around -40C. Though in this instance my intention was to make N2O3 as pure as possible, in future trials I think I’m going to feed some pure O2 into the ampoules containing N2O3 before sealing to convert it all to NO2/N2O4.
Another idea that came to mind was possibly adding weak H2O2 to the Sulfuric (pure sulfuric doesn’t easily react with pure nitrite so it could be used instead of water to dilute) and maybe that could slowly decompose producing more NO2. I know there’s other ways of directly producing NO2 but RFNA plus most metals doesn’t give much NO2 per gram and RFNA plus sulfur gives impurities in the condensed liquid.

649315E6-1F91-441B-B0DF-091B0DE0B45F.jpeg - 3.5MB

caterpillar - 19-4-2021 at 01:43

RFNA ignites turpentine- I made such an experiment. No need for N2O4.

DBX Labs - 11-6-2021 at 13:46

I uploaded a video on the subject of synthesizing and ampouling nitrogen oxides through the route of nitrite decomposition earlier today.

Heres the link:
https://youtu.be/pcvypNVYJYc

MineMan - 14-6-2021 at 17:17

Did you dissolve a green sour patch kid in it??